1. Kinetic studies of keto-enol and other tautomeric equilibria by flash photolysis.
Jakob Wirz. Department of Chemistry. University of Basel. Klingelbergstrasse ...
1
Kinetic studies of keto-enol and other tautomeric equilibria by flash photolysis Jakob Wirz Department of Chemistry University of Basel Klingelbergstrasse 80, CH-4056 Basel, Switzerland 1
INTRODUCTION
1
2
METHODS
3
2.1
Flash Photolysis
3
2.2
Derivation of the Rate Law for Keto–Enol Equilibration
3
2.3
Halogen Titration Method
8
2.4
pH–Rate Profiles
9
2.5
General Acid and General Base Catalysis
11
3
EXAMPLES
13
4
RATE–EQUILIBRIUM RELATIONSHIPS
21
4.1
The Brønsted Relation, Statistical Factors, and the Acidity of Solvent-Derived Species
⊕
(H and H2O)
21
4.2
Mechanism of the "Uncatalyzed" Reaction
24
4.3
The Marcus Model of Proton Transfer
25
5
CONCLUSION AND OUTLOOK
29
6
REFERENCES
29
1 Introduction Erlenmeyer was first to consider enols as hypothetical primary intermediates in a paper published in 1880 on the dehydration of glycols.1 Ketones are inert towards electrophilic reagents, in contrast to their highly reactive enol tautomers. However, the equilibrium concentrations of simple enols are generally quite low. That of 2-propenol, for example, amounts to only a few ppb in aqueous solutions
2 of acetone. Nevertheless, many important reactions of ketones proceed via the more reactive enols, and enolization is then generally rate-determining. Such a mechanism was put forth in 1905 by Lapworth,2 who showed that the bromination rate of acetone in aqueous acid was independent of bromine concentration and concluded that the reaction is initiated by acid-catalyzed enolization, followed by fast trapping of the enol by bromine (Scheme 1). This was the first time that a mechanistic hypothesis was put forth on the basis of an observed rate law. More recent work has shown that the reaction of bromine with various acetophenone enols in aqueous solution takes place at nearly, but not quite, diffusion-controlled rates.3 O
slow H2O/H
OH
fast Br2
O Br
+ HBr
Scheme 1
In 1978, we observed that flash photolysis of butyrophenone produced acetophenone enol as a transient intermediate, which allowed us to determine the acidity constant KaE of the enol from the pH–rate profile (section 2.4) of its decay in aqueous base.4 That work was a sideline of studies aimed at the characterization of biradical intermediates in Norrish Type II reactions and we had no intentions to pursue it any further. Enter Jerry Kresge, who had previously determined the ketonization kinetics of several enols using fast thermal methods for their generation. He immediately realized the potential of the photochemical approach to study keto–enol equilibria and quickly convinced us that this technique should be further exploited. We were more than happy to follow suit and to cooperate with this distinguished, inspiring and enthusing chemist and his cherished wife Yvonne Chiang, who sadly passed away last year. Over the years, this collaboration developed into an intimate friendship of our families. The present chapter is an account of what has been achieved. Several reviews in this area appeared in the years up to 1998.5-10 The enol tautomers of many ketones and aldehydes, carboxylic acids, esters and amides, ketenes, as well as the keto tautomers of phenols have since all been generated by flash photolysis to determine the pH–rate profiles for keto-enol interconversion. Equilibrium constants of enolization, KE, were determined accurately as the ratio of the rate constants of enolization, kE, and of ketonization, kK, Equation 1. KE = kE/kK Equation 1. Kinetic determination of equilibrium constants of enolization
Strong bases in dry solvents are usually used in organic synthesis to generate reactive enol anions from ketones. Nevertheless, the kinetic studies discussed here were mostly performed on aqueous solutions. Apart from the relevance of this medium for biochemical reactions and green chemistry, it has the advantage
3 of a well-defined pH-scale permitting quantitative studies of acid and base catalysis.
2 Methods 2.1 Flash Photolysis The technique of flash photolysis, introduced in 1949 by Norrish and Porter,11 now covers time scales ranging from a few femtoseconds to seconds and has become a ubiquitous tool to study reactive intermediates. Most commonly, light induced changes in UV-Vis optical absorption are monitored, either at a single wavelength (kinetic mode) or spectrographically at a given delay with respect to the light pulse used for excitation (spectrographic mode, pump–probe spectroscopy). Instruments of a conventional design,12 which employ an electric discharge to produce a strong light flash of sub-millisecond duration, usually have sufficient time resolution and are then most suitable to study the kinetics of keto–enol tautomerization reactions. Nowadays, instruments using a Q-switched laser as an excitation source having durations of a few nanoseconds (laser flash photolysis) are much more widespread. These techniques are well-known, and their properties, pitfalls and limitations have been described.13-15
2.2 Derivation of the Rate Law for Keto–Enol Equilibration Activation energies for unimolecular 1,3-hydrogen shifts connecting ketones and enols are prohibitive, so that thermodynamically unstable enols can survive indefinitely in the gas phase or in dry, aprotic solvents. Ketones are weak carbon acids and oxygen bases, enols are oxygen acids and carbon bases. In aqueous solution, keto–enol tautomerization proceeds by proton transfer involving solvent water. In the absence of buffers, three reaction pathways compete, as shown in Scheme 2. OH KaK O K k0'K+ kH'K cH
K KE
rate determining E
kHK cH + k0K
OH E
'E k0'E+ kOH cOH
O
E
k0 + kOH cOH
KaE
E Scheme 2. Acid-, base-, and “uncatalyzed” reaction paths of keto–enol tautomerism.
Four species participate in the tautomerization reaction, the ketone (K, e.g., acetone), the protonated ketone (K⊕), the enol (E), and its anion (E ). These species are connected through two thermodynamic cycles. The Gibbs free
4 energies for the individual elementary reactions r of any cycle must add up to naught, Equation 2. Σ∆rG° = 2.3RTΣpKr = 0 Equation 2
For the cycle K → E → E + H⊕ → K we get pKE + pKaE – pKaK = 0, where KE is the equilibrium constant of enolization and KaE and KaK are the acidity constants of E and K, respectively; KaK is defined in the direction opposite to the last process of the cycle so that pKaK must be subtracted. Similarly, the equilibrium constant for carbon deprotonation of the protonated ketone, K⊕ → E + H⊕, can be replaced by pKE + pKaK⊕, where pKaK⊕is the acidity constant of K⊕. Thus, the equilibrium properties of Scheme 2 are fully defined by the three equilibrium constants KE, KaE, and KaK⊕. We turn to the kinetic parameters. When an enol E is rapidly generated in a concentration cE(t = 0) exceeding its equilibrium concentration cE(∞), the decrease of cE(t) may be followed in time by, for example, some absorbance change as in flash photolysis. Deprotonation or protonation of carbon atoms is generally slow relative to the equilibration of oxygen acids with their conjugate bases. Therefore, carbon acids and bases have been called pseudo-acids and pseudo-bases. Proton transfer reactions involving carbon are the rate-determining elementary steps of the tautomerization reactions. A shaded oblique line is drawn across these reactions in Scheme 2. Thus we posit that the protonation equilibria on oxygen that are associated with the ionization constants KaE and KaK⊕ are established at all times during the much slower tautomerization reactions. This assumption leads to a pH-dependent first-order rate law for keto–enol tautomerization reactions, Equation 14, that will be derived below and is found to hold in general. The pre-equilibrium assumption adopted for oxygen acids is, thereby, amply justified. We define equilibrium constants as concentration quotients, as in Equation 3 for KaE and KaK⊕. Provided that the experiments are done at low and constant ionic strengths, I ≤ 0.1 M, these can be converted to thermodynamic constants, Ka°, using known or estimated activity coefficients.16 KaE = cE (t) cH⊕/cE(t) and KaK⊕ = cK(t) cH⊕/cK⊕ (t) Equation 3
The total concentration of the enol and its anion is cE,tot(t) ≡ cE(t) + cE (t); inserting Equation 3 we can express the concentrations cE and cE as a function of proton concentration cH⊕, Equation 4. cH ! K aE cE t = E c t and cE! t = E c t K a + cH ! E,tot K a + cH ! E,tot
()
()
()
()
Equation 4
Protons and hydroxyl ions are not consumed by the reaction K E. A temporary shift in the relative concentrations of K and E may, however, lead to a
5 change in proton concentration cH⊕ due to rapid equilibration with K⊕ and E , respectively. To avoid this complication, the conditions are generally chosen such that cH⊕ remains essentially constant during the reaction by using either a large excess of acid or base, or by the addition of buffers in near neutral solutions (pH = 7 ± 4). However, the addition of buffers usually accelerates the rates of tautomerization. We first consider reactions taking place in wholly aqueous solutions, that is, in the absence of buffers. The handling of rate constants obtained with buffered solutions will be discussed in section 2.5. To derive the general rate law for keto–enol equilibration, we consider each of the rate-determining elementary reaction steps shown in Scheme 2 separately, beginning with enol ketonization reactions. The relevant rate constants for the rate-determining ketonization reactions are kH⊕K and k0K for C-protonation of E by H⊕ and solvent water, respectively, and kH⊕'K and k0'K for C-protonation of E by H⊕ and water (Scheme 2). We use primed symbols k' for the rate constants referring to ketonization of the anion E . As we shall see in a moment (Equation 7), the terms k0K and kH⊕'KKaE are both independent of pH and may be combined to a single term kucK. The associated, seemingly "uncatalyzed" reactions are therefore kinetically indistinguishable and additional information is required to determine, which of the corresponding mechanisms is the dominant one (see section 4.2). We assume that the rate-determining reactions shown in Scheme 2 are elementary reactions, so that the corresponding rate laws are equal to the product of a rate constant and the concentrations of the reacting species. a) Acid-catalyzed ketonization: The rate for ketone formation by carbon protonation of the enol E is given by Equation 5, where the right-hand expression is obtained by substituting cE(t) using Equation 4. c ! vHK! = kHK! cH ! cE t = kHK! cH ! E H c t K a + cH ! E,tot
()
()
Equation 5
b) Base-catalyzed ketonization: Pre-equilibrium ionization of E generates the more reactive anion E , which may be protonated on carbon by the general acid water in the rate-determining step, Equation 6. For pH-values well below pKaE, the concentration cH⊕ is much greater than KaE, so that it may be neglected in the denominator of Equation 6. The rate of this reaction is then inversely proportional to cH⊕, i.e., proportional to cOH . This "apparent" base catalysis saturates at pH-values above pKaE, when E is converted to E . The concentration cH⊕ then becomes much smaller than KaE and may be neglected in the denominator of Equation 6. K aE K 'K 'K vOH! = k0 cE! t = k0 c t K aE + cH ! E,tot
()
Equation 6
()
6 c) "Uncatalyzed" ketonization. At pH-values near neutral, a pH-independent rate of ketonization is frequently observed, which may be attributed to several different mechanisms (see section 4.2): carbon protonation of E by water or a concerted transfer of the enol proton to carbon through one or more solvent molecules, and carbon protonation of E by the proton, Equation 7. For pH
